A Single-Component Liquid-Phase Hydrogen Storage Material

Nov 9, 2011 - We describe a liquid-phase hydrogen storage material that is a liquid under ambient ... and Facile Metal-Free Hydrogen Transfer Reaction...
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A Single-Component Liquid-Phase Hydrogen Storage Material Wei Luo, Patrick G. Campbell, Lev N. Zakharov, and Shih-Yuan Liu* Department of Chemistry, University of Oregon, Eugene, Oregon 97403-1253, United States

bS Supporting Information ABSTRACT: The current state-of-the-art for hydrogen storage is compressed H2 at 700 bar. The development of a liquid-phase hydrogen storage material has the potential to take advantage of the existing liquid-based distribution infrastructure. We describe a liquid-phase hydrogen storage material that is a liquid under ambient conditions (i.e., at 20 °C and 1 atm pressure), air- and moisture-stable, and recyclable; releases H2 controllably and cleanly at temperatures below or at the proton exchange membrane fuel cell waste-heat temperature of 80 °C; utilizes catalysts that are cheap and abundant for H2 desorption; features reasonable gravimetric and volumetric storage capacity; and does not undergo a phase change upon H2 desorption.

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afe, efficient storage and delivery of hydrogen is essential for the development of a hydrogen-based energy infrastructure. Storage of hydrogen as a compressed gas (up to 10 000 psi/700 bar) is the current state-of-the-art,1,2 but to increase the storage density and mitigate the risks associated with storage and transport of high-pressure gas, numerous condensed-phase hydrogen storage approaches are currently under investigation. These include metal hydrides,3 sorbent materials,4 and chemical hydride5 systems.6 Boron- and nitrogen-containing chemical hydrides have attracted much attention because of their high gravimetric hydrogen densities and favorable kinetics of hydrogen release.7,8 Ammonia borane (H3N BH3, AB), with a gravimetric density of 19.6 wt % H2, is one of the most promising candidates among the chemical hydride materials. AB has both hydridic and protic hydrogens, facilitating H2 release under mild conditions. However, while the release of H2 from AB and its derivatives has been extensively investigated,9,10 AB is a solid material that releases H2 at its melting point and cannot serve as liquid fuel without dilution (e.g., with a solvent), which necessarily reduces its hydrogen storage capacity.11 The appeal of a safe, liquid-phase hydrogen storage material is clear.12 The U.S. has a network of over 150 000 miles (244 000 km) of pipeline dedicated to delivering liquid petroleum products, and many nations worldwide have similar networks in place.13 The transition to a hydrogen-based energy economy will be greatly facilitated if it can take advantage of the existing liquidbased distribution channels such as pipelines, tankers, and retail outlets. Two potential liquid-phase hydrogen storage materials that have received recent attention in the literature are formic acid, HCO2H,14 and hydrous hydrazine, N2H4 3 H2O.15 One disadvantage of these compounds is that they have decomposition pathways that potentially generate side products that are toxic to fuel cell catalysts (e.g., CO and NH3) in addition to safety r 2011 American Chemical Society

Scheme 1. A Single-Component Liquid-Phase Hydrogen Storage Material

concerns (e.g., for hydrazine16). Liquid organic hydrides (i.e., hydrocarbons) are another class of potential hydrogen carriers, but for carbon-rich systems, the hydrogen liberation step is strongly endothermic, typically requiring reaction temperatures of 350 500 °C,17 well above the “waste heat” temperature of 80 90 °C provided by a standard proton exchange membrane (PEM) fuel cell. This limitation can be overcome somewhat by the incorporation of heteroatoms into the carbon scaffold. Pez, Scott, and Chang of Air Products Corporation have studied the use of 9-ethylcarbazole as a hydrogen storage material and demonstrated the dehydrogenation of 9-perhydroethylcarbazole at 150 200 °C in a series of patents.18,19 In 2010, Tsang et al. published a method to regenerate spent 9-ethylcarbazole fuel using molecular H2 and an alumina-supported ruthenium catalyst.20 One drawback to 9-ethylcarbazole hydrogen storage is that the spent fuel material is a solid at temperatures up to 60 °C. Research in our lab focuses on the fundamental chemistry and applications of boron nitrogen heterocycles (i.e., cyclic molecules containing carbon, boron, and nitrogen).21 27 We recently reported the synthesis of 1,2-BN-cyclohexane,28 a cyclic amine borane. Cyclic amine boranes have not been as extensively explored as AB and its acyclic derivatives in H2 storage applications, arguably because of the lack of synthetic access. We noted that 1,2-BN-cyclohexane (mp: 63 °C) cleanly releases hydrogen to form a trimer when thermally activated. As part of an ongoing program directed toward hydrogen storage by BN heterocycles,29,30 we disclose herein the development of BN-methylcyclopentane (1) (Scheme 1), which is an air- and moisture-stable liquid at room temperature. We report that 1 is capable of releasing 2 equiv of H2 per molecule of 1 (4.7 wt %) both thermally, at temperatures above 150 °C, and catalytically using a variety of cheap and abundant metal halides, at temperatures Received: September 19, 2011 Published: November 09, 2011 19326

dx.doi.org/10.1021/ja208834v | J. Am. Chem. Soc. 2011, 133, 19326–19329

Journal of the American Chemical Society

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Scheme 2. Synthesis of Charged Liquid Fuel 1

below 80 °C. The exclusive product of dehydrogenation is the trimer 2, which is also a liquid at room temperature. We also demonstrate that the conversion of the spent fuel 2 back to the charged fuel 1 can be accomplished in high yield under relatively mild conditions, making this system a viable candidate for liquidphase hydrogen storage in mobile and carrier applications. The synthesis of 1 is illustrated in Scheme 2. Treatment of the bis-N-protected amine 3 with neat BH3 3 Et3N at 160 °C for 48 h generated heterocycle 4, which was not isolated. The crude mixture was diluted with THF, and then KH and HF 3 pyridine were added to generate the charged fuel 1. The product was purified by column chromatography under ambient conditions (i.e., in the presence of oxygen and moisture), and 1 was isolated in 51% overall yield from 3. Compound 1 is a liquid at room temperature with a melting point of 18 °C. At lower temperatures, we were able to grow crystals of 1 that were suitable for single-crystal X-ray diffraction (XRD) analysis, thus unambiguously confirming our structural assignment (see the ORTEP in Scheme 2). We determined that heterocycle 1 is thermally stable at 35 °C as a neat liquid [see the Supporting Information (SI) for details]. However, upon heating at 150 °C for 1 h in the absence of solvent, each of three molecules of 1 releases 2 equiv of H2 in forming the trimer 2 (eq 1), which is also a liquid at room temperature (mp: 9 °C). Thus, the hydrogen desorption from charged fuel 1 to form the spent fuel 2 does not involve a phase change, a critical property for a liquid-phase H2 carrier in terms of actual application in fuel cells.11,31 Additionally, 1 and 2 display low volatility, with boiling points of approximately 335 and 378 °C, respectively (determined at reduced pressure and extrapolated to 1 atm), reducing the risk of introducing potential fuel cell poisons into the H2 stream.

Metal-catalyzed dehydrogenation of AB has attracted growing attention from the perspective of hydrogen storage. Although many important advances have been made with Pd, Ru, Ir, and

Figure 1. Automated volumetric buret measurements of H2 release catalyzed by metal chloride complexes.

other noble-metal catalysts,9 the development of convenient-tohandle, cheap, abundant, and efficient catalysts with low toxicity is of considerable interest. Our research has focused on first-row transition-metal halide catalysts.32 To our knowledge, there have been only a limited number of publications discussing the dehydrogenation AB with these metal salts. Ramachandran and co-workers reported the use of NiCl2 and CoCl2 as catalysts for methanolysis of AB,33 and Jagirdar et al. used CoCl2, NiCl2, and CuCl2 as reactants to assist in the hydrolysis of AB.34 The use of iron halide salts for AB dehydrogenation in the solid state has recently been reported by Chen et al.35 To find the most effective metal halide catalyst for the dehydrogenation of 1, we screened anhydrous F, Cl, Br, and I complexes of Fe, Co, Ni, and Cu at 5 mol % catalyst loading in toluene.36 We monitored the reaction progress by 11B NMR spectroscopy. The postulated dimer intermediate 5 (Table S1 in the SI) was visible in the 11B NMR spectrum (96 MHz, toluene; 3.0 ppm, doublet, 1JBH = 117 Hz) but could not be isolated.37 41 The conversions after 5 min at 80 °C are listed in Table S1. We found that in general, bromide complexes are the most reactive toward formation of 2, followed by chloride and then iodide complexes, and that fluoride complexes are almost completely inactive. Copper, nickel, and cobalt halides are more reactive than iron halides. The two most active catalysts in this study were NiBr2 and CuBr, both of which achieved 76% conversion to 2 in 5 min. All of the selected chloride, bromide, and iodide complexes can completely dehydrogenate 1 to release 2 equiv of H2 per molecule of 1 in less than 30 min. The presence of a catalyst was essential for H2 desorption at 80 °C. No H2 release was observed after 1 h at 80 °C without a catalyst. To understand further the differences between various iron, cobalt, nickel, and copper chloride complexes, we performed several dehydrogenation experiments using an automated gas buret apparatus.42 Chloride complexes were chosen for this study because they are significantly cheaper than bromide complexes. In a general procedure, we dissolved 75 mg of 1 in toluene with 5 mol % catalyst and submerged the reaction flask in an 80 °C oil bath. Varying the metal resulted in markedly different hydrogen release profiles (Figure 1). CoCl2 promoted the release of 2 equiv of H2 from 1 in ca. 7 min, and the CuCl2- and NiCl2-catalyzed reactions were complete in under 10 min. The reactions with iron complexes were slower; both FeCl3 and FeCl2 promoted the 19327

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Scheme 3. Conversion of Spent Fuel 2 Back to Charged Fuel 1

Figure 2. Large-scale dehydrogenation of 1 using 5 mol % FeCl2 without solvent.

Figure 3. (a) H2 evolution trace for sequential addition of charged fuel 1. (b) H2 evolution trace with Hg poisoning of the catalyst.

desorption of 2 equiv of H2 in ca. 15 min. Interestingly, for the cobalt- and nickel-catalyzed reactions, the initial rate of H2 desorption (i.e., from t = 0 to t for 1.0 equiv of H2) is apparently lower than the rate over the time period from 1.0 to 2.0 equiv. We repeated the automated buret measurement experiments illustrated in Figure 1 at 50 °C and noted that the time required for complete H2 desorption exceeded 4 h for all catalysts (see the SI for details). This suggests that the reaction temperature has a significant effect on the rate of dehydrogenation.

To demonstrate the potential utility of our material as a simple-to-operate, single-component liquid system, we performed a large-scale dehydrogenation of 1 (10 mmol, the maximum capacity of our buret apparatus) without additional solvent using 5 mol % FeCl2 as a catalyst.43 Figure 2 shows that 2 equiv of H2 were released from the neat material in ca. 20 min at 80 °C. At the conclusion of the reaction, the spent fuel 2 was isolated in 95% yield. Noteworthy is the induction period of ca. 4 min before significant H2 release was observed.44 The FeCl2 catalyst changed form during the course of the reaction, becoming an amorphous black powder. At 5 mol % loading, the catalyst could be reused without loss of activity (Figure 3a). Adding fresh fuel 1 to the reaction flask after dehydrogenation of the old batch had completed resulted in immediate H2 desorption from the new batch of fuel (tested up to three times, the capacity of our buret apparatus; see the SI for details). We also performed dehydrogenation experiments with reduced catalyst loading. With 3 mol % FeCl2, the rate of dehydrogenation decreased significantly; 2 equiv of H2 were released in 40 min. With 1 mol % FeCl2, only 0.4 equiv of H2 were released after 70 min. To probe the nature of the active catalyst species, Hg was added to the reaction mixture, and a reduced rate of hydrogen release was observed, consistent with heterogeneous catalysis (Figure 3a vs 3b).45 Recyclability is crucial to the success of any hydrogen storage system. For H2 desorption of AB, a variety of monomeric (e.g., cyclotriborazane, cyclopentaborazane, and borazine) and polymeric (e.g., polyamino- and -iminoboranes and polyborazylene) spent fuel products can be produced depending on the dehydrogenation conditions, thus making this system less well-defined and arguably potentially more challenging to regenerate.46 Recently, Sutton and Gordon demonstrated that one spent fuel product of AB dehydrogenation, polyborazylene, can be regenerated with hydrazine in liquid ammonia.47 49 The hydrogen desorption of storage material 1 to form 2 is a clean process. We believe that the well-defined molecular nature of the spent fuel 2 should facilitate the development of a regeneration process. We determined that treatment of 2 with methanol for 12 h at room temperature produces the bismethoxy species 6 (Scheme 3), as confirmed by single-crystal XRD analysis.50 Subsequent treatment of 6 with LiAlH4 afforded the charged fuel 1 in 92% overall yield. This “regeneration” sequence was performed using the product of our 10 mmol-scale dehydrogenation experiment (Figure 2) to demonstrate the recyclability of our hydrogen storage system on a larger scale. We acknowledge that the described regeneration scheme is not yet optimal from an energetic point of view (i.e., the use of highly energetic LiAlH4 ultimately needs to be avoided). In summary, we have developed an air- and moisture-stable liquidphase hydrogen storage material, 1, that does not undergo a phase 19328

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Journal of the American Chemical Society change upon H2 desorption. We have discovered a series of first-row transition-metal halide catalysts that are capable of releasing 2 equiv of H2 from 1 in less than 30 min in toluene at 80 °C at modest catalyst loadings. We have demonstrated that neat liquid 1 can quantitatively release H2 in the presence of FeCl2 catalyst. Furthermore, we have shown that the spent fuel 2 can be converted back to the charged fuel 1 in good yield. The disclosed hydrogen storage material could represent a viable H2 storage option for mobile, portable, and carrier applications. Current efforts are being directed toward developing a more cost- and energy-efficient regeneration procedure.

’ ASSOCIATED CONTENT

bS

Supporting Information. Experimental procedures, spectroscopic data, and crystallographic data (CIF). This material is available free of charge via the Internet at http://pubs.acs.org.

’ AUTHOR INFORMATION Corresponding Author

[email protected]

’ ACKNOWLEDGMENT This paper is dedicated to Prof. Daniel G. Nocera. S.-Y.L. gratefully acknowledges the U.S. Department of Energy (DEFG36-08GO18143) for financial support. Funding for the University of Oregon Chemistry Research and Instrumentation Services was furnished in part by the NSF (CHE 0923589). ’ REFERENCES (1) Eberle, U.; Felderhoff, M.; Sch€uth, F. Angew. Chem., Int. Ed. 2009, 48, 6608. (2) Jorgensen, S. W. Curr. Opin. Solid State Mater. Sci. 2011, 15, 39. (3) Graetz, J. Chem. Soc. Rev. 2009, 38, 73. (4) Zhao, D.; Yuan, D.; Zhou, H.-C. Energy Environ. Sci. 2008, 1, 222. (5) Stephens, F. H.; Pons, V.; Baker, R. T. Dalton Trans. 2007, 2613. (6) McWhorter, S.; Read, C.; Ordaz, G.; Stetson, N. Curr. Opin. Solid State Mater. Sci. 2011, 15, 29. (7) Hamilton, C. W.; Baker, R. T.; Staubitz, A.; Manners, I. Chem. Soc. Rev. 2009, 38, 279. (8) Marder, T. B. Angew. Chem., Int. Ed. 2007, 46, 8116. (9) For an overview, see: Staubitz, A.; Robertson, A. P.; Manners, I. Chem. Rev. 2010, 110, 4079. (10) Bowden, M.; Autrey, T. Curr. Opin. Solid State Mater. Sci. 2011, 15, 73. (11) For an example of the use of liquid-phase amine borane fuel blends, see: Mal, S. S.; Stephens, F. H.; Baker, R. T. Chem. Commun. 2011, 47, 2922. (12) For an overview, see: Teichmann, D.; Arlt, W.; Wasserscheid, P.; Freymann, R. Energy Environ. Sci. 2011, 4, 2767. (13) CIA World Factbook. https://www.cia.gov/library/publications/ the-world-factbook/ (accessed Aug 31, 2011). (14) For an overview of HCO2H as a potential energy carrier, see: Enthaler, S.; von Langermann, J.; Schmidt, T. Energy Environ. Sci. 2010, 3, 1207. (15) Singh, S. K.; Xu, Q. J. Am. Chem. Soc. 2009, 131, 18032. (16) For a leading reference, see: Choudhary, G.; Hansen, H. Chemosphere 1998, 37, 801. (17) Coughlan, B.; Keane, M. A. Catal. Lett. 1990, 5, 89. (18) Pez, G. P.; Scott, A. R.; Cooper, A. C.; Cheng, H. U.S. Patent 7,101,530, 2006. (19) Pez, G. P.; Scott, A. R.; Cooper, A. C.; Cheng, H.; Wilhel, F. C.; Abdourazak, A. U.S. Patent 7,351,395, 2008. (20) Eblagon, K. M.; Rentsch, D.; Friedrichs, O.; Remhof, A.; Zuettel, A.; Ramirez-Cuesta, A. J.; Tsang, S. C. Int. J. Hydrogen Energy 2010, 35, 11609.

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(21) Marwitz, A. J. V.; Abbey, E. R.; Jenkins, J. T.; Zakharov, L. N.; Liu, S.-Y. Org. Lett. 2007, 9, 4905. (22) Abbey, E. R.; Zakharov, L. N.; Liu, S.-Y. J. Am. Chem. Soc. 2008, 130, 7250. (23) Marwitz, A. J. V.; Matus, M. H.; Zakharov, L. N.; Dixon, D. A.; Liu, S.-Y. Angew. Chem., Int. Ed. 2009, 48, 973. (24) Campbell, P. G.; Abbey, E. R.; Neier, D.; Grant, D. J.; Dixon, D. A.; Liu, S.-Y. J. Am. Chem. Soc. 2010, 132, 18048. (25) Abbey, E. R.; Zakharov, L. N.; Liu, S.-Y. J. Am. Chem. Soc. 2010, 132, 16340. (26) Lamm, A. N.; Garner, E. B.; Dixon, D. A.; Liu, S.-Y. Angew. Chem., Int. Ed. 2011, 50, 8157. (27) Abbey, E. R.; Zakharov, L. N.; Liu, S.-Y. J. Am. Chem. Soc. 2011, 133, 11508. (28) Luo, W.; Zakharov, L. N.; Liu, S.-Y. J. Am. Chem. Soc. 2011, 133, 13006. (29) Campbell, P. G.; Zakharov, L. N.; Grant, D. J.; Dixon, D. A.; Liu, S.-Y. J. Am. Chem. Soc. 2010, 132, 3289. (30) Matus, M. H.; Liu, S.-Y.; Dixon, D. A. J. Phys. Chem. A 2010, 114, 2644. (31) Jiang, H.-L.; Singh, S. K.; Yan, J.-M.; Zhang, X.-B.; Xu, Q. ChemSusChem 2010, 3, 541. (32) Catalysis without Precious Metals; Bullock, R. M., Ed.; WileyVCH: Weinheim, Germany, 2010. (33) Ramachandran, P. V.; Gagare, P. D. Inorg. Chem. 2007, 46, 7810. (34) Kalidindi, S. B.; Indirani, M.; Jagirdar, B. R. Inorg. Chem. 2008, 47, 7424. (35) He, T.; Wang, J.; Wu, G.; Kim, H.; Proffen, T.; Wu, A.; Li, W.; Liu, T.; Xiong, Z.; Wu, C.; Chu, H.; Guo, J.; Autrey, T.; Zhang, T.; Chen, P. Chem.—Eur. J. 2010, 16, 12814. (36) See Table S1 for details. The Lewis acid catalysts AlCl3 and ZnCl2 were also screened; they were found to give incomplete dehydrogenation and intractable product mixtures and were not pursued further. (37) Shaw, W. J.; Linehan, J. C.; Szymczak, N. K.; Heldebrant, D. J.; Yonker, C.; Camaioni, D. M.; Baker, R. T.; Autrey, T. Angew. Chem., Int. Ed. 2008, 47, 7493. (38) Jaska, C. A.; Temple, K.; Lough, A. J.; Manners, I. J. Am. Chem. Soc. 2003, 125, 9424. (39) Chen, Y.; Fulton, J. L.; Linehan, J. C.; Autrey, T. J. Am. Chem. Soc. 2005, 127, 3254. (40) Fulton, J. L.; Linehan, J. C.; Autrey, T.; Balasubramanian, M.; Chen, Y.; Szymczak, N. J. Am. Chem. Soc. 2007, 129, 11936. (41) Rousseau, R.; Schenter, G. K.; Fulton, J. L.; Linehan, J. C.; Engelhard, M. H.; Autrey, T. J. Am. Chem. Soc. 2009, 131, 10516. (42) Zheng, F.; Rassat, S. D.; Helderandt, D. J.; Caldwell, D. D.; Aardahl, C. L.; Autrey, T.; Linehan, J. C.; Rappe, K. G. Rev. Sci. Instrum. 2008, 79, No. 084103. (43) The approximate bulk price of FeCl2 (ca. $0.30 kg 1) is based on a survey of international suppliers. Also see: http://www.alibaba. com/product-gs/296803887/Ferrous_Chloride_4_Hydrate.html (accessed Sept 2, 2011). (44) No induction period before hydrogen release was observed for experiments run in solution (see Figure 1). (45) Jaska, C. A.; Manners, I. J. Am. Chem. Soc. 2004, 126, 9776. (46) Wright, W. R.; Berkeley, E. R.; Alden, L. R.; Baker, R. T.; Sneddon, L. G. Chem. Commun. 2011, 47, 3177. (47) Sutton, A. D.; Burrell, A. K.; Dixon, D. A.; Garner, E. B.; Gordon, J. C.; Nakagawa, T.; Ott, K. C.; Robinson, P.; Vasiliu, M. Science 2011, 331, 1426. (48) For an alternative regeneration scheme for AB spent fuel, see: Davis, B. L.; Dixon, D. A.; Garner, E. B.; Gordon, J. C.; Matus, M. H.; Scott, B.; Stephens, F. H. Angew. Chem., Int. Ed. 2009, 48, 6812. (49) Sutton, A. D.; Davis, B. L.; Bhattacharyya, K. X.; Ellis, B. D.; Gordon, J. C.; Power, P. P. Chem. Commun. 2010, 46, 148. (50) See the SI for details. Correspondence concerning X-ray crystallography should be directed to Lev Zakharov (e-mail: lev@ uoregon.edu). 19329

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